Question for ceg4048

[Jose]

Hi everyone,

This is an extract of the article on CO2 by ceg4048,

"The relationship in practical terms therefore is that if Hobbyist "A" has tap water measuring ph 7.2 and kH 10 ( high levels of carbonate and bicarbonates) then 30 ppm of dissolved CO2 may only cause his pH to drop to 7.0

Conversely, Hobbyist "B" has tap water also measuring 7.2 but kH 6. More acid can form in his water (because of less carbonate and bicarbonate levels), so 30 ppm dissolved CO2 in his water will result in a pH drop to 6.8."

Ok Id like to understand this better because Im a bit confused. I thought that to get 30 ppms (or any ppm of CO2) the change in pH would have to be the same (e.g 1(aprox) for 30 ppm) no matter what KH the water is. This is what I understood from Tom Barr.

So to understand this better:
Hobbiest A, water of KH 10, His pH would drop 1 unit.
Hobbiest B, water of KH 6, His pH would drop 1 unit as well.

To get another ppm of CO2 the water would change X units of pH, but it should be the same X amount for a different water with another kh.

We can check this with the pH-KH table. You can try doing it with two different kh waters. The result is that the change in pH is the same for both if we have same co2 ppm. I know ph kh tables dont work very well for real waters but for this comparison its appropriate, I would've thought..

What am I missing?

Thanks all!!

 

[markk]

What am I missing?

Hi Jose,

I won't even attempt to give you the specific science, but bear in mind that the key issue is that water with a higher kH has more buffering capacity - hence more CO2 is required for a given pH change.

I'm not sure which ph/kH tables you've been looking at but the numbers you've quoted above seem to tie in with these charts:-

http://www.barrreport.com/forum/barr-report/co2-enrichment/11862-co2-ph-kh-table
http://www.practicalfishkeeping.co.uk/content.php?sid=5264#
http://www.gpodio.com/co2_chart.asp

(if anything, the pH drop for the water with 6 kH is a little understated - in theory!)

regards

Mark

 

[Jose]

(if anything, the pH drop for the water with 6 kH is a little understated - in theory!)

Thanks for replying Markk,
No nevermind the exact numbers. Its all about the idea. Ive read those Tom Barr posts but Ill read through them again see if I can find exactly what I mean (unless I misunderstood)

 

[Jose]

I won't even attempt to give you the specific science, but bear in mind that the key issue is that water with a higher kH has more buffering capacity - hence more CO2 is required for a given pH change.

More CO2 is needed but the X amount of CO2 dissolved changes pH in the same Y ammount no matter what hardness the water (theoretically)

 

[Rahms]

More CO2 is needed but the X amount of CO2 dissolved changes pH in the same Y ammount no matter what hardness the water (theoretically)

I'm not really familiar with the science here, but looking at the tables I can say this statement is dangerously on the verge of being wrong . The change in ppm is not linear. Going from 3ppm to 30ppm (27ppm increase) is always going to drop your pH by 1, yes. However, dropping your pH by 2 will not take you to 57 ppm (i.e. 3ppm+27ppm+27ppm), it will take you to 300ppm. The example in your first post has different starting ppms on that table.

It seems to imply "hobbyist A" has tapwater coming out at 18.9ppm CO2 though! I think it's saying that until he doses above 18.9ppm he will simply be displacing something else (carbonates?), so there will be no pH change even though his dissolved CO2 is increasing. That or its an unrealistic example and his water would be coming out at closer to 8 pH!

 

[Jose]

All starting points are from equilibrium with air.

 

[Rahms]

All starting points are from equilibrium with air.

Not really sure where this fits in with what I said. It's just... a fact.

 

[EnderUK]

Sorry miss understood.

You're thinking of pH as a linear scale. It's not really, it's a log scale.

 

[markk]

More CO2 is needed but the X amount of CO2 dissolved changes pH in the same Y ammount no matter what hardness the water (theoretically)

Jose,

I don't believe that is true - although I'll concede I may be either wrong or just not understanding what you're trying to prove/understand.

Apologies if any of this is pitched at the wrong level (either to technical or too simplistic) - I don't know what your background is.

Firstly, any CO2 that isn't dissolved is irrelevant to pH. Yes - there will be an equilibrium between gaseous and aqueous CO2 but I'm assuming that we're talking about the ppm of dissolved CO2.

Secondly, the pH scale is not linear. pH is broadly speaking a measurement of the concentration of H+ ions and the amount of H+ ions needed to drive any given solution from, say, pH 7 to pH 6 is very different from the amount needed to move it from pH 6 to pH 5. Not strictly relevant to your comment above - but it's an important point, particularly when we start comparing movement in CO2 ppm (linear scale) and pH (logarithmic scale).

and thirdly, is buffering again But getting more technical (I had to resort to some revision) - have a look at the equation (and text) on this page:-

http://mbrewer.edublogs.org/2012/03/07/the-carbonatebicarbonate-buffer-system/

CO2 will tend to form carbonic acid in water, which will in turn tend to split into an H+ ion and bicarbonate and the bicarbonate will tend to further split into another H+ ion and carbonate.

For our purposes we pump CO2 in at the left and on the right we get some H+ ions that drive our pH value (but remember the differences in scale). With no other factors in play (lets say it's pure water), there will be a direct relationship between CO2 in and pH 'out' - but the specific pH value is dependant on the balance of the different equilibria in the equation referenced.

But if you now add some Sodium bicarbonate or Potassium bicarbonate, these will raise your kH and in aqueous solution will disassociate and add bicarbonate and carbonate to your solution - i.e. 2 of the same molecules that are already at play in the CO2 equilibria above. These will push the balance of that equilibria to the left - hence reducing the number of H+ ions and increasing the pH. The more (bi)carbonate you add, the more you will shift the balance.

Approaching it from the other direction, if I start with a solution of 'high' kH, I have to add much more CO2 on the left to achieve the same number H+ ions on the right (i.e. lower the pH) - as the relatively high levels of (bi)carbonates will tend to 'absorb' them.

Hopefully I've got that right - and not gone off on a complete tangent! (but the revision was useful to me at least)

Cheers

Mark

 

[Jose]

Sorry guys but hopefully someone will answer my question since it hasnt been by now. Ill give a clearer example hopefully. So guys grab a ph kh table.

Tank number 1: kh=10
If we take water from it and let it equlibrate with the atmosphere in a glass it will give a ph reading of around 8.3 and the co2 ppm will be around 1.5 (its just an example dont get hung up on numbers please) and this is just water + carbonates. All data is extracted from the table.
Now we inject co2 until we have 30 ppm and the ph should drop to around 7. This is a 1.3 units ph drop.

Tank number 2: kh=15
We do the same, put it in a glass and let it reach equikibrium. Ph should read around 8.5 with 1.5ppm of co2 again.
Now we add co2 once more until 30 ppm and ph should read around 7.2 (from chart) This is again a ph drop of 1.3 units.

In both cases(kh 10 and kh 15) ph dropped 1.3 units to get 30 ppm co2 dissolved. So we need the same ph drop independently of kh( hardness) to get to a certain co2 ppm.

This are all numbers extracted from the ph kh chart. The only suposision I made is the co2 ppm for equilibrium. This might not be 1.5 ppm of co2 but its irrelevant what the exact number is really.

 

[Rahms]

but thats completely different to the original example you posted, ofcourse no-one has answered it.

I acknowledged in an earlier post going 1.5ppm -->30ppm will create the same pH drop in the table. But in the original example, the water shows ~11ppm and ~19ppm from the tap. So the issue is whether the real-life measurement is accurate at all levels of CO2, or only beyond a certain point, which is what markk was explaining when talking about buffering.

 

[markk]

Sorry guys but hopefully someone will answer my question since it hasnt been by now. Ill give a clearer example hopefully. So guys grab a ph kh table

OK - sorry for the confusion - I see where you're coming from now.

Going back to your original question, ceg4048 just assumed 2 water samples with the same starting pH but different kHs.

In your examples you've assumed 2 water samples with the same starting CO2 but different kHs.

You're both then adjusting the CO2 to 30ppm. In ceg4048's case, the change in CO2 is different for each sample (he made no assumptions about CO2 levels at the beginning) but in your case it's the same for each sample.

Just two different uses of the same table - does that help?

regards

Mark

 

[Jose]

Yeap I think you guys got it right this time. Well done.

But still You would think that you have to let tap water rest and get to equilibrium to measure pH. This way you can measure the real pH change due to your CO2 injection without the error of the tap CO2. Wouldnt you?
If this is as you say guys then the example is not about kh but about different CO2 concentratiosn from the tap which doesnt make sense.
He is trying to make a point about waters with different kh(I wouldve thought). In one of them you need to change pH more than in the other one. But this is not backed up by the table or is it?
Unless the equilibrium CO2 ppm is different for each water (I dont think its that different).

So, ceg's starting points arent necessarily from equilibrium with the atmosphere? Different kh and same ph just means there is something else in the water other than carbonates. Could be CO2 or not. If its not CO2 then the pH drop should be approx the same (right?).

Cheers!

 

[markk]

markk puts his Devil's Advocate hat on

But still You would think that you have to let tap water rest and get to equilibrium to measure pH.

Why? Th pH is the pH - but see below.

This way you can measure the real pH change due to your CO2 injection without the error of the tap CO2. Wouldnt you?

But you're injecting CO2 into your tank. The real pH change - the one that is important - is in your tank water. The starting point is your tank water parameters before lights on/CO2 on and the end point is the pH when you reach your desired CO2 levels.

The kH/pH chart is telling you that for higher kH, it gets progressively harder to push the pH below around 7 and for lower kH you can drop below 7 and beyond very quickly if you're not careful.

If, for your water, the pH swing is too high for your livestock, then you need a rethink...

If this is as you say guys then the example is not about kh but about different CO2 concentratiosn from the tap which doesnt make sense.

The example is about kH, pH and CO2 - it just doesn't care about the starting CO2 levels (or whether it is tap or tank water).

He is trying to make a point about waters with different kh(I wouldve thought). In one of them you need to change pH more than in the other one. But this is not backed up by the table or is it?
Unless the equilibrium CO2 ppm is different for each water (I dont think its that different).

Well - it could be. I can't find anything to support this but if you go back to all of those carbonate buffering equilibria, if you have higher kH then you are pushing those equilibria towards the left (CO2 + H2O). This should have an effect on the equilibrium between 'atmospheric' CO2 and dissolved CO2.

So, ceg's starting points arent necessarily from equilibrium with the atmosphere? Different kh and same ph just means there is something else in the water other than carbonates. Could be CO2 or not. If its not CO2 then the pH drop should be approx the same (right?).

I think they are, as you say, just examples and they are just demonstrating the general concept. They tie in with the 'theory' behind the tables, as do the examples you gave. You just need to remember that they don't care about the starting point of CO2.

regards

Mark

 

[Jose]

Why? Th pH is the pH - but see below.

Because this is the only way to get a consistent measurement since co2 in tap water will vary. The idea is to be able to just measure a pH drop when you are fine tuning your CO2. No need to measure kH.

But you're injecting CO2 into your tank. The real pH change - the one that is important - is in your tank water. The starting point is your tank water parameters before lights on/CO2 on and the end point is the pH when you reach your desired CO2 levels.

Exactly the best way to get a starting point of your tank water is from the equilibrium with the atmosphere (glass of water) which will always be the same(aprox). This way you can rule out measuring kh which is a huge source of error.

I think we got to the bottom of it. His starting points are not from the equilibrium with atmosphere but still the question stands if he is comparing khs. pH will need the same jump to get to x ppm of CO2 for different kh waters. This is of course if the starting point is the same one (e.g equilibrium w/o co2 injection). In other words we need our pH to change the same amount no matter what kh., if everything else is the same (starting and finishing CO2 ppms). So pH jump (not CO2 volume injected) that we need is independent of water hardness.

 

[Jose]

Here are some references to the 1 unit ph change to get 30 ppm CO2. Notice there is no reference to kh.

http://www.barrreport.com/forum/barr-report/co2-enrichment/217799-co2-impact-of-kh-and-kh

here are his exact words from another one:

"Say you have 30ppm CO2 in a KH of 10 and a 30ppm CO2 in a KH of 1.

The pH difference is? 1.0 pH units.

6.0 vs 7.0 pH.

Without adding CO2?
About 7.2 and 8.2"

And the link:
http://www.barrreport.com/forum/bar...1102-different-co2-forms-under-low-vs-high-ph

So....cegs words seem to be at least misleading.

 

[markk]

Here are some references to the 1 unit ph change to get 30 ppm CO2. Notice there is no reference to kh.

http://www.barrreport.com/forum/barr-report/co2-enrichment/217799-co2-impact-of-kh-and-kh

here are his exact words from another one:

"Say you have 30ppm CO2 in a KH of 10 and a 30ppm CO2 in a KH of 1.

The pH difference is? 1.0 pH units.

6.0 vs 7.0 pH.

Without adding CO2?
About 7.2 and 8.2"

And the link:
http://www.barrreport.com/forum/bar...1102-different-co2-forms-under-low-vs-high-ph

So....cegs words seem to be at least misleading.

This specific quote

The pH difference is? 1.0 pH units.

is referring to the difference in pH between two specific, different kH solutions, both with 30ppm CO2.

The 1.0 is pure coincidence (pick a different kH and check).

The point he is making is that if you take away the extra CO2, and implicitly let it drop back to some sort of 'atmospheric' equilibrium, both solutions will change pH by about 1.2 i.e. the same- which is what I thought we were talking about?

 

[Jose]

The 1.0 is pure coincidence (pick a different kH and check).

I have done it. Its always around 1 unit pH. Its not coincidence. You arent reading the table properly then.

 

[Jose]

The point he is making is that if you take away the extra CO2, and implicitly let it drop back to some sort of 'atmospheric' equilibrium, both solutions will change pH by about 1.2 i.e. the same- which is what I thought we were talking about?

Exactly this is what we are talking about. Are we reading same things?

 

[markk]

I have done it. Its always around 1 unit pH. Its not coincidence. You arent reading the table properly then.

for CO2 @ 30ppm

- kH 1: pH = 6.0
- kH 4: pH = 6.6

Have I got that wrong?

or do I need some fresh air